Sciencemadness Discussion Board - Butyric acid from butanol & KMnO4

Butyric acid from butanol & KMnO4

Hockeydemon - 19-7-2013 at 17:34

I was trying to make some butyric acid by oxidizing butanol with KMnO4, but I have had no such luck.

I figured I could simply add a hot solution KMnO4 to butanol, and the stopping point would become apparent by the smell. Because of this I didn't really take note of how much KMnO4 I added in total, but I started with 100mL of butanol, and I added an absurd amount of KMnO4 solution relative to the 100mL of butanol.

A reaction was clearly taking place because of how exothermic the addition of KMnO4 was. But no matter how much I seemed to add I never really got anywhere.

It had crossed my mind that I allowed the reaction to get too hot which caused the butanol to boil off, but at the same time in order for that heat to be generated in the first place wouldn't some of the butanol had to have been oxidized to butanoic acid?

The KMnO4 is a very pretty purple color in solution, and when added to the clear butanol a brown sludge is formed.

On a side note I took about 2mL of that sludge, and added ~1mL of nitric acid to it. The mixture instantly started to bubble, then went clear & formed 2 layers. Any clue as to what took place in that reaction?

bfesser - 19-7-2013 at 17:57

Which isomer of <a href="https://en.wikipedia.org/wiki/Butanol#Isomers" target="_blank">butanol</a> <img src="../scipics/_wiki.png" /> do you have? What were the approximate temperature and concentration of your <a href="http://en.wikipedia.org/wiki/Potassium_permanganate" target="_blank">potassium permanganate</a> <img src="../scipics/_wiki.png" /> (KMnO4

solution?

In the reaction, the KMnO4 would be reduced to <a href="http://en.wikipedia.org/wiki/Manganese(IV)_oxide" target="_blank">manganese(IV) oxide</a> <img src="../scipics/_wiki.png" /> (MnO2

, this is your brown sludge.

[Edited on 7/20/13 by bfesser]

Hockeydemon - 19-7-2013 at 18:19

I figured I'd need to have provided more detail to get any real answers. I guess I was more looking for X was wrong in your process.

I'll have to redo it to give you the conc of the KMnO4 (any recommendation?). The temp was at least 100C since it was at a boil. It is 1-butanol.

The solution containing the butanol & KMnO4 was very hot, and steaming quite a bit.

bfesser - 19-7-2013 at 18:51

I'd recommend a little lit. research before you proceed. <a href="http://library.sciencemadness.org/library/books/vogel_practical_ochem_3.pdf" target="_blank">Vogel's Bible</a> <img src="../scipics/_pdf.png" /> has a few KMnO4 oxidations (including one of isobutanol).

[Edited on 7/20/13 by bfesser]

Hockeydemon - 19-7-2013 at 18:59

I know I'll be reading up a bit more before I give it another go - it just appeared to be really straight forward when I read about it initially.

I was wondering though.. After I form butanoic acid I should have a solution of that + MnO2 (& H2O & K). Could I distill off the butanoic acid, and then use nitric acid to nitrate the MnO2 to form manganese nitrate + H2O? Since they are immiscible I could separate the two, and be left with manganese nitrate?

Or does HNO3 + MnO2 not form manganese nitrate since you edited that part out of your post?

Lambda-Eyde - 19-7-2013 at 19:34

Quote: Originally posted by Hockeydemon

I was wondering though.. After I form butanoic acid I should have a solution of that + MnO2 (& H2O & K). Could I distill off the butanoic acid, [...]

Obviously not, since butyric acid has a boiling point that is 63 degrees higher than that of water. If you had checked Wikipedia (for instance) in advance, you would have known that. You also would have known that alkaline permanganate oxidizes butyric acid to carbon dioxide. I don't know if this is what happened here (seemingly in neutral conditions), but it might be a possibility. Dichromate would anyways be a better oxidizing agent.

What sort of apparatus are you doing this in? An open beaker? How do you think you're going to tackle the overpowering smell of 100+ g of hot butyric acid as the reaction proceeds? I'd do it in a two- or three-neck flask, alternately a flask with a claisen adapter with an attached reflux condenser leading to a trap filled with a lye solution, and an addition funnel in the remaining neck. After adding all the oxidant and letting the reaction come to the halt, I'd probably distill off the water, then the acid, and finally dry and distill again. All done in an apparatus sealed with an alkaline lock, of course (and don't forget the suck-back traps).

But these are just my personal ramblings and gut-feel. Look up a tried and true published procedure.

Quote: Originally posted by Hockeydemon

Or does HNO3 + MnO2 not form manganese nitrate since you edited that part out of your post?

No, manganese dioxide contains Mn(IV), while manganese nitrate contains Mn(II). You'd have to reduce the dioxide first. There is a prep for preparing manganese(II) sulfate from manganese(IV) dioxide using sulfuric acid and hydrogen peroxide. I do not know if you can do the same with nitric acid.

Your check-list for next time:

Gather all relevant physical data such as boiling points, molecular weights, concentrations etc.
Look up a published procedure for this procedure or an analogous one (e.g. the oxidation of 1-propanol to propanoic acid) and follow it. Write out your steps. Know what to do in order to handle smells and eventual spills.
Consider scaling down from your original scale. You've just wasted 100 mL of butanol and a considerable amount of permanganate. Had you done it at 1/10 of the scale you would have wasted less reagents. Also, 10 mL of butyric acid is a whole lot less unpleasant to spill than 100 mL of that beast.
Don't you ever again think that organic chemistry is easy or straight forward. It never is

Magpie - 20-7-2013 at 08:22

KMnO4 is a strong oxidant and can indeed be used to chew up organics and spit them out as CO2.

I suggest you try a milder oxidant - say nitric acid. Here's a procedure for conversion of an alcohol to an acid which may prove beneficial:

http://www.sciencemadness.org/talk/viewthread.php?tid=13122

Runaway reactions, even explosions, are a risk when mixing an alcohol (a fuel) with a strong oxidant. I think the best technique is to get the oxidant up to reaction temperature and then very gingerly and slowly add the alcohol to keep the reaction going. When all the alcohol has been added and the reaction seems to have stopped, then apply some heat to assure maximum completeness of the reaction.

[Edited on 20-7-2013 by Magpie]

Acidum - 20-7-2013 at 12:33

First thing - KMnO₄ is very powerful oxidizer, and it also has several levels of oxidizing depending on reaction conditions. Depending on pH, concentrations and temperature, it can do things from mild oxidations to formations of carboxylic acids to total oxidation of organic compounds to CO₂. So, check on conditions BEFORE trying to synthetize something. Most likely You created vast amounts of CO₂...

Oh, and definitely this:

Quote: Originally posted by Lambda-Eyde

Your check-list for next time:

Gather all relevant physical data such as boiling points, molecular weights, concentrations etc.
Look up a published procedure for this procedure or an analogous one (e.g. the oxidation of 1-propanol to propanoic acid) and follow it. Write out your steps. Know what to do in order to handle smells and eventual spills.
Consider scaling down from your original scale. You've just wasted 100 mL of butanol and a considerable amount of permanganate. Had you done it at 1/10 of the scale you would have wasted less reagents. Also, 10 mL of butyric acid is a whole lot less unpleasant to spill than 100 mL of that beast.
Don't you ever again think that organic chemistry is easy or straight forward. It never is

Hockeydemon - 22-7-2013 at 01:58

Firstly - thank you for all of the helpful responses. I'm just now starting O-chem so my apologies on my naivety.

Quote: Originally posted by Lambda-Eyde

What sort of apparatus are you doing this in? An open beaker? How do you think you're going to tackle the overpowering smell of 100+ g of hot butyric acid as the reaction proceeds? I'd do it in a two- or three-neck flask, alternately a flask with a claisen adapter with an attached reflux condenser leading to a trap filled with a lye solution, and an addition funnel in the remaining neck. After adding all the oxidant and letting the reaction come to the halt, I'd probably distill off the water, then the acid, and finally dry and distill again. All done in an apparatus sealed with an alkaline lock, of course (and don't forget the suck-back traps).

I was just doing it in a beaker yes. I had an aspirator next to me that I was going to throw on once I smelled anything foul.

I have a 3-neck flask that I can setup with an addition funnel, & reflux condenser. Why would my reflux condenser lead to a solution of lye? I understand the premise of neutralizing the acid, but the acid should never leave the reflux condenser? I'm also unfamiliar with an 'alkaline lock' & wasn't able to turn anything up readily with google.

Quote: Originally posted by Magpie

KMnO4 is a strong oxidant and can indeed be used to chew up organics and spit them out as CO2.

I suggest you try a milder oxidant - say nitric acid. I think the best technique is to get the oxidant up to reaction temperature and then very gingerly and slowly add the alcohol to keep the reaction going. When all the alcohol has been added and the reaction seems to have stopped, then apply some heat to assure maximum completeness of the reaction.

Adding the alcohol to the oxidant would only be viable with HNO3 correct? If I have the KMnO4 in excess and add the alcohol to it wouldn't it proceed to 'chew up' the alcohol producing CO2?

-----

Can someone correct anything I get wrong in the following.. Forgive me if this is a vast over simplification.

The butanol is first oxidized to the aldehyde butanal with the addition of permanganate because the electron rich single bonded oxygen atom cleaves off two hydrogen from the butanol leaving butanal & H2O. This changes the oxidation state of the Mn to +6 (from +7).

Why is the aldehyde able to pull off another oxygen atom from the manganese trioxide?

-----
I looked up the isobutyric acid synthesis in Vogel's book (pg. 668 exp:5.122). His uses fairly large quantities, of everything & ends up with a yield of only 45g isobutyric acid which isn't very much. I can of course scale his synthesis down (I would have to - I don't have a 5L rb flask). But I would like to use the product to make other aromatic compounds.

-----
If I used nitric acid as the oxidizer would each HNO3 molecule only be able to perform one oxidation? In other words an oxygen molecule would be liberated from the HNO3 to form butanal, and the HNO2 would simply decompose (2HNO2 -> H2O + NO + NO2)? Then another HNO3 molecule would give up an oxygen atom to form the carboxylic acid?

bfesser - 22-7-2013 at 05:01

It is clear from the context what you mean by "aromatic compound," but in the future, be careful with that term, as it's a synonym for <a href="http://en.wikipedia.org/wiki/Aromaticity" target="_blank">arenes</a> <img src="../scipics/_wiki.png" />.

Also, when Lambda-Eyde wrote "alkaline lock," what he meant (archaically) was the gas trap which will neutralize acid vapor ("a trap filled with a lye solution").

woelen - 22-7-2013 at 05:22

May I suggest that next time you do the reactions first on a small scale. It is a pity that you spoilt 100 ml of 1-butanol and an insane amount of KMnO4. This failure must have cost you quite some money. Bad bad bad.

Next time take 2 ml or so in a test tube and add a little amount of oxidizer. This saves you a lot of money and it strongly reduces the risk of severe fire and/or explosion.

Another thing is that KMnO4 as solid oxidizer, mixed simply with an organic compound does not yield nice results. You get a horrible black/brown slurrie and you rip apart most of the organic molecules leaving you with useless crap and a lot of CO2 and H2O.

I'm quite sure that you can get much better results if you dissolve 1-butanol in dilute sulphuric acid and then add KMnO4. In sufficiently acidic aqueous solution the KMnO4 will be reduced all the way down to Mn(2+) instead of MnO2 and the reaction is much more controlled. The alcohol-group first is oxidized to an aldehyde-group and after that it is oxidized to the acid. So, you'll first get CH3CH2CH2C(=O)H and then you'll get CH3CH2CH2COOH. There is no ripping apart of the C-skeleton under these conditions.

Try this in small quantities in a test tube and see whether you obtain the smell of the acid before you continue with larger amounts. Isolating the acid from the mix will require a decent distillation setup. You first boil off water, then the acid and you'll be left with excess H2SO4 and MnSO4/K2SO4 mix. Do not boil off to dryness, because that will lead to charring of the butyric acid.

Magpie - 22-7-2013 at 11:37

Quote: Originally posted by Hockeydemon

Adding the alcohol to the oxidant would only be viable with HNO3 correct? If I have the KMnO4 in excess and add the alcohol to it wouldn't it proceed to 'chew up' the alcohol producing CO2?

Not necessarily. This is likely concentration dependent.

At this point, where you are exploring to find what works at a small scale, I wouldn't worry too much about stoichiometry. I would base my initial testing on what has worked for others in similar alcohol oxidations. In the oxidation of 1,3 propanediol to malonic acid I assumed the reaction shown was appropriate stoichiometry. But the actual ratio of HNO3/OH that I used was twice that shown in my equation. My choice was totally based on what worked in the reference by Seymour where he was making oxalic acid.

As Lambda-Eyde says, in organic chemistry what works is often not terribly predictable.

Hockeydemon - 22-7-2013 at 19:25

Okay I went ahead a tried again small scale in 3 test tubes. Each test tube contained 4mL of 1-butanol.

In test tube 1 I had 2mL of 18M H2SO4, and proceeded to slowly add 4mL of HNO3 with frequent stirring. I allow this to happen in the open atmosphere which allowed the reaction to get fairly hot but still touchable. This resulted in a loss of liquid in the end. I ended up with ~3mL of a dark black liquid that had a slightly sweet syrupy smell to it. It also of course evolved a lot of NO & NO2 during the addition. Thoughts?

In test tube 2 I had 2mL of 18M H2SO4 along with the addition of less than 1g of KMnO4. I heated this with a Bunsen burner for a little bit with frequent stirring. Not too much happened other than dark purple/gray liquid formed. I'm guessing since there was a bit of dark purple remaining that I did not allow the reaction to complete?

In test tube 3 I only added HNO3, and I kept the test tube in cold tap water so the reaction did not get past luke warm. After the first 2mL of HNO3 was added two layers formed. The top layer was a dark green color, and the bottom layer was a light green color. After ~4mL of HNO3 was added it was all the light green color, but a reaction was still occurring because it would gradually heat up if I took it out of the water & vigorous stirring would increase the evolution of NO & NO2 gas. After ~7mL of HNO3 was added the solution was a light yellow color, and it would continue to heat up if not in the water bath, stirring still exacerbated gas evolution. I then placed a stirbar into the test tube, and while in the water bath allowed it to vigorously stir until I could no longer get a gas to evolve. I then took it out of the water, and allowed the stirring to continue & I applied heat to the test tube. This caused quite a bit of gas to evolve, but I noticed no real chance to the solution even after I allowed it to cool back down to room temperature.

None of these variations or steps produced anything that smelled unpleasant. Any thoughts or suggestions? I am curious what the various intermediates I formed in test tube 3 are if I did not form butyric acid. Is the green the formation of the aldehyde?

Magpie - 22-7-2013 at 19:51

Try this:

Place 3-4mL of 6M HNO3 in a test tube. Add to this 1-2 drops of your butanol. Heat up carefully using a hot water bath until you can see a vigorous reaction taking place. See if it smells of butyric acid.

woelen - 23-7-2013 at 05:18

I see you don't use any water at all. All of these reactions are a form of destuctive oxidation and do not yield what you want. The black stuff you obtain sometimes is called 'gunk' or 'goop'

and is totally useless. The sweet smell may be due to broken down molecules having nitro-groups or nitrate groups attached to them. Beware, some of these compounds are very toxic and strong carcinogens.

Try the following:

Take 3 ml of water and add 1 ml of conc. H2SO4. This will cause the liquid to heat up considerably. Then add 0.5 ml of 1-butanol. Next add a small spatula of KMnO4 and swirl gently. Assure that you have excess 1-butanol. It may take quite some time and quite some shaking/stirring before all of the KmnO4 has dissolved. Smell the liquid. I'm quite sure that this reaction leads to formation of butyric acid. If it indeed works, then the next step is to work towards a mechanism with precisely determined stoichiometry, the first step I described here is checking whether the reaction works.

[Edited on 23-7-13 by woelen]

Magpie - 23-7-2013 at 08:46

Quote: Originally posted by Hockeydemon

I looked up the isobutyric acid synthesis in Vogel's book (pg. 668 exp:5.122). His uses fairly large quantities, of everything & ends up with a yield of only 45g isobutyric acid which isn't very much. I can of course scale his synthesis down (I would have to - I don't have a 5L rb flask). But I would like to use the product to make other aromatic compounds.

I looked up the preparation of isobutyric acid in Vogel in the forum library. It gives a preparation using alkaline permanganate that gives a 92.5% yield. What more could you want. Scale it down if it is too large a batch.

BTW I wouldn't waste any more effort on using nitric acid. Not when you have a procedure provided in detail in Vogels.

Waffles SS - 23-7-2013 at 10:10

Nitric acid usually use for oxidation Diol alcohol and i think this is not common method for oxidation Primary alcohol to Carboxylic acid but i like to try it for oxidation 1-Butanol next week and i will share the result.

I dont suggest Alkaline/Permanganate oxidation method because:
-You should use huge amount of water for making only small amount of carboxylic acid(you need 100lit of water for 1 kg of Carboxylic acid !)
-You should cool huge amount of water and also reduce it by vacuum for best performance
-This method take lot of time(~12hours)

Nitric acid oxidation will be great if it can oxidize primary alcohol to carboxylic acid safe and fast

@Magpie,
Have you ever tried this method for primary alcohol?

[Edited on 23-7-2013 by Waffles SS]

Magpie - 23-7-2013 at 11:57

No, waffles, malonic acid is the only acid I have made using nitric acid. As you say, it was made from a diol. The procedure is convenient because the workup is so simple.

Brewster gives a procedure for converting sucrose to oxalic acid using nitric acid. I suppose sucrose could be considered a polyol.

Hockeydemon - 24-7-2013 at 21:09

Okay I performed another round of experiments with multiple variations with no luck thus far.

Trial 1: I did exactly as Magpie suggested with the HNO3. I place 4mL of 6M HNO3 in a test tube. Then I added 4 drops of 1-butanol and suspended this in a flask of water. I then heated the flask of water with a Bunsen burner for awhile. The water was just about to boil (small bubbles forming on the inside of the flask) when the solution in the test tube abruptly started to boil vigorously. I then removed the heat source and allow the reaction to continue. The reaction continued for about 15 minutes (new bubbles formed). During the reaction oxides were released. The end result was a clear liquid with a faint NO2 smell though it is a reasonable assumption this could be just remnant of the gas evolved.

Trial 2: I did what Woelen said (same measurements & .5g KMnO4 in each) with two variations. The first variation I allowed to react with no cooling which allowed the reaction to get quite hot. The solution went very dark at first, and after several minutes it formed two layers. The top layer was roughly .5mL of a tan liquid. The bottom layer was clear, and made up the majority of the solution.

The second variation I did the whole reaction at 5C in an ice water bath. After about 10 minutes of vigorous stirring in the ice bath I removed the solution. The solution was a very dark color - I'd say a dark brown. This also formed two layers with the top one being comprised of roughly .5mL of the same tan liquid. The bottom one was of course the dark brown liquid.

Neither of these produced any unpleasant smells. I also made sure to pull a small amount of the bottom layer out with a pipette, and I placed that on a paper towel to see if it smelled any different. I did not notice any difference. Since nothing occurred in either test tube variation I decided to heat both solutions with a torch (gradually). Once the solutions both got up to temperature a reaction took place that made each solution go clear. The first variation test tube (not cooled) was crystal clear with a small dark ring of residue around the top of the liquid (this is possibly just KMnO4 dust that was on the walls of the test tube). The second variation (cooled test tube) also went clear, but had a slight tan tint to it.

Both liquids smelled quite pleasant. I can't think of anything that it reminds me of, but I did think it smelled good enough that I was upset I was limited to wafting the odor towards my face rather than actually smelling it. I wish I knew what it was (I know you said possibly toxic molecules with nitrate groups attached).

----

On top of those I also did a 1% scale of Vogel's isobutryic acid synthesis (obviously using 1-butanol instead). I placed 1mL of 1-butanol, and ~.1-.2g (non-mg scale) of sodium bicarbonate into 1.5mL of water. I placed this into a 100mL round bottom flask which I suspended in a 1L beaker of iced salt water over a stir-plate. I then added 1.4g of KMnO4 to 27.5mL of water and mixed until dissolved. I then added this to the rxn vessel which was at 5C. The instructions say to continue stirring for 3-4hrs, and then allow the rxn to come to room temperature over the next 12 hours. I will however be leaving for work shortly so the rxn will continue stirring for the next 8-9 hours. I will stop it when I get home, and it will sit for another 7 hours while I sleep. I will then continue with Vogel's instructions from there & let you know of the results.

[Edited on 25-7-2013 by Hockeydemon]

woelen - 24-7-2013 at 23:28

Did you use dilute sulphuric acid and added 1-butanol to that, and then added KMnO4? This is as I described it.

I am surprised that you did not get any butyric acid with my method. I have the impression that you obtain an ester (n-butyl butyrate), but this would surprise me, given all the water present in the mix.

If you indeed obtain an ester, then the reaction does work, but immediately after formation of the butyric acid it reacts with excess 1-butanol to obtain the ester. This, however, is not something which I would expect, but it is the only explanation which I have for what you observe. Especially because you obtain a clear liquid in the end and not some black gunk.

In the experiments with KMnO4 as oxidizer you can smell safely, you do not have to worry about very toxic compounds in the mix. The formation of the strongly carcinogenic compounds only occurs when there is destructive oxidation with HNO3 (if you get plumes of NO2, together with white fumes, then you'd better leave the place for a while and allow the fumes to diffuse away).

Hockeydemon - 24-7-2013 at 23:55

I added H2O to the test tube, then H2SO4, then butanol, and then while stirring added the KMnO4.

After reading the scent profile of n-butyl butyrate I would say that is a fair description of what I smelled. Am I correct in assuming that the top .5mL of the solution would be the butyl butyrate ester, and the bottom layer was mostly water?

Since the butyric acid is reacting with excess 1-butanol should I try running the experiment again, and half the 1-butanol quantity?

woelen - 25-7-2013 at 00:10

I do not think that the top 0.5 ml is pure n-butyl butyrate, but it most likely contains quite some of this ester. I expect it to be a mix of the ester and 1-butanol.

Maybe you could get better results when you use less 1-butanol, but you still have the complication that the ester is formed. This is something which complicates things a lot. I myself have experience with ethanol being oxidized to acetic acid and there, the problem of formation of ethyl acetate in the aqueous solutions is nearly non-existent. I expected it to be the case with 1-butanol as well, but apparently things are otherwise.

I am afraid that an organic chemist must jump in with good ideas.

DutchChemistryBox - 25-7-2013 at 02:59

Maybe you can try to make the Butyl butanoate on purpose with the description of woelen but then with less water.
And seperate the ester from the rest and then add natriumhydroxide to make Sodium butanoate.
After that you can evaporate the solution so you'll have an almost alcohol and water free solid.

If you distill this with concentrated H2SO4 I would say that your product is butanoic acid.

Or am I thinking way out of the box?

Hockeydemon - 25-7-2013 at 18:16

I've just allowed the mixture described in Vogel's book to sit for 12+ hours, and I have a few questions going forward with the instructions.

It says to then filter off the manganese dioxide, and concentrate the filtrate under reduced pressure then cool.

This is simply saying distill the filtrate under reduced pressure right? Does what pressure I get to matter?

It then says cover the solution with a layer of ether, and acidify with dilute sulfuric acid.

I don't have ether or any non-polar solvents currently. I have several aprotic, and protic polar solvents. I guess I could make some ether - I do have some pure ethanol in the freezer. Is this the only way around this?

It also doesn't give a quantity for what acidifying the solution requires. Is this something that I'm supposed to know?

adamsium - 25-7-2013 at 21:04

Concentrating the solution under reduced pressure means to remove some of the solvent and possibly excess reagents under vacuum, leaving a more concentrated solution of the product. This is what you would often do on a rotary evaporator (though not to dryness in this case, it seems).

It's probably sufficient to acidify until the solution is just acidic.

Magpie - 25-7-2013 at 21:40

Can you smell any butyric acid?

Hockeydemon - 25-7-2013 at 21:49

No :/. It just smells like a solution of KMnO4.. I followed the procedure from Vogel exactly - it is a very easy procedure to follow.

Magpie - 25-7-2013 at 21:57

IMO if you can't smell it at this point it probably doesn't exist.

But if you feel it's worth the effort to follow through on the procedure then that is up to you.

Hockeydemon - 25-7-2013 at 22:09

Is it possible there was steric interference preventing the reaction from occurring that wasn't there with the iso form of the molecule? Or just some reason that the iso form would work, and the 1-butanol would not be able to in this case specifically?

[Edited on 26-7-2013 by Hockeydemon]

adamsium - 25-7-2013 at 22:50

I found a patent that you might find interesting. They used nitric acid (and lots of it) and supposedly achieved very god yields. You'll notice that the alcohol was added slowly, rather than all at the beginning.

Attachment: GB771583A.pdf (224kB)
This file has been downloaded 680 times

Rich_Insane - 26-7-2013 at 15:30

Would the Jones reagent not work in this case? I thought it was quite suitable for oxidation of alcohols into carboxylic acids....

Hockeydemon - 26-7-2013 at 16:01

So I had 25mL of nitric acid left, and I went ahead and did reaction IV in the document adamsium provided. I scaled it down to 25mL HNO3, 1.5mL H2O, & 4.8mL of 1-butanol. Unfortunately I wasn't able to maintain 25-30C well enough, and a run away reaction happened. It was a neat occurrence actually.. The temperature wasn't too far off 30C, and with the addition of 1 drop it just took off, and billowed out an amazing amount of nitrogen dioxide. It made a waterfall of dark red smoke from my work bench to the floor, and crawled across the floor.

Anyways I'm going to make another batch of nitric acid, and give the rxn another go. I'll probably try to keep the water bath ~15C in hopes that will keep the reaction around 30C with vigorous stirring.

Thanks for the document adamsium. Hopefully I'll have time tomorrow to try again.

Hockeydemon - 28-7-2013 at 18:11

Okay, I went ahead and redid the reaction same way as in my previous post, but I made sure to keep the reaction at 30C. I ended up with a light green liquid I still don't smell any butyric acid however.

I'm not sure why I don't smell any butyric acid :/. I'm also not sure what I'm making if it isn't making butyric acid..

Edit:
I read here that HNO3 + alcohol --> nitrate ester + H2O, and I have found no mention (outside of that patent linked above) that I can use HNO3 to oxidize the alcohol to an acid.

Is this correct?

[Edited on 29-7-2013 by Hockeydemon]

Waffles SS - 29-7-2013 at 09:08

Usually Nitrate ester prepared by action of conc. nitric acid(>90%) on alcohol + sulfuric acid as catalyst in below 15c(above this temp run away occur and maybe explosion happen)
I think you made Butyl Butyrate.Try to hydrolysis it by NaOH solution and then add Acid for making butyric acid.

[Edited on 29-7-2013 by Waffles SS]

adamsium - 29-7-2013 at 09:21

I hadn't heard of using nitric acid, either. Based on what I've learned in my organic chem classes so far, I'd have done something similar to what you were doing. I'm quite curious about this, too.

Magpie - 29-7-2013 at 09:40

Welcome to organic chemistry. There are many general procedures for syntheses but the yields in specific cases can vary widely. I am surprised that the procedure in Vogel didn't work, but then again why would it call out the isobutyric form if there wasn't some specificity involved. I have seen a procedure specifically for n-butyric acid in Gatterman, 1937 edition (forum library) but it uses a different precursor, ie, ethyl malonic acid, IIRC.

I know you want to make this acid using the precursor you have on hand, ie, n-butanol. I am also very puzzled as to why you have had no success so far. There are two other oxidants that you might try, ie, potassium (or sodium) dichromate w/H2SO4, and sodium (or potassium) persulfate.

If you really, really wanted to make butyric acid you could try other routes, eg, Grignard reaction using n-propanol and CO2 (dry ice), or hydrolysis of propyl cyanide. I have made an acid using the Grignard route but not the cyanide route. You likely don't want to go to these extremes. I, too, have to believe it is possible to make it from n-butanol. If you find a route you will have something worthwhile to contribute to SM.

[Edited on 29-7-2013 by Magpie]

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Here's a synthesis of butyl butyrate that provides some insight into making butyric acid. It looks like control of mole ratios and temperature are important. As indicated, if there is too much alcohol the ester is made, as previously pointed out upthread by woelen:

http://www.orgsyn.org/orgsyn/default.asp?formgroup=basenpe_f...

One could approach the making of n-butyric acid by first making the aldehyde and isolating it. That should be easily oxidized to the acid, possibly using just stirring in air with a little heat and time.

[Edited on 30-7-2013 by Magpie]

[Edited on 30-7-2013 by Magpie]

Hockeydemon - 29-7-2013 at 18:08

Quote: Originally posted by Waffles SS

I think you made Butyl Butyrate.Try to hydrolysis it by NaOH solution and then add Acid for making butyric acid.

I know I'm making butyl butyrate in the test tube when I follow the instructions Woelen gave me. The liquid is clear, insoluble in water, and smells pleasant.

I don't know that in the reaction I followed with nitric acid I made butyl butyrate. The liquid was a light green/light cyan color, and there was no other layer in the liquid. I added NaOH solution to until until the pH went up to 7 which made the liquid yellow. Then I added a little more, and the solution went to orange (pH of ~9). The solution was still miscible, and there was plenty of water in the mixture. I then added H2SO4 to the mixture, and it simply returned to it's original form (light green/light cyan).

I will however make that ester of butyl butyrate using Woelen's instructions, and separate the ester. Then I will hydrolyze the ester, and re-acidify it.

Quote: Originally posted by Magpie

Here's a synthesis of butyl butyrate that provides some insight into making butyric acid. It looks like control of mole ratios and temperature are important. As indicated, if there is too much alcohol the ester is made, as previously pointed out upthread by woelen:

http://www.orgsyn.org/orgsyn/default.asp?formgroup=basenpe_f...

One could approach the making of n-butyric acid by first making the aldehyde and isolating it. That should be easily oxidized to the acid, possibly using just stirring in air with a little heat and time.

The link you provided just takes me to the search engine for the website. Can you tell me what you searched for, and I can look it up?

Magpie - 29-7-2013 at 18:16

Quote: Originally posted by Hockeydemon

The link you provided just takes me to the search engine for the website. Can you tell me what you searched for, and I can look it up?

search for: n-butyl n-butyrate

Hockeydemon - 30-7-2013 at 04:21

Since I can't seem to oxidize 1-butanol to butyric acid with KMnO4 I think I'm going to try as you suggested Magpie & isolate the aldehyde. Rather than using KMnO4 at all I am going to try a copper catalyzed dehydrogenation to it's aldehyde. From there I should hopefully be able to carboxylize it.

We shall see how it goes I guess - I'll give it a go later in the afternoon.

How should I go about converting the aldehyde to the carboxylic acid?

[Edited on 31-7-2013 by Hockeydemon]

Waffles SS - 30-7-2013 at 05:45

Quote: Originally posted by Hockeydemon

How should I go about converting the aldehyde to the carboxylic acid?

Quote: Originally posted by Magpie

It should be easily oxidized to the acid, possibly using just stirring in air with a little heat and time

Butyraldehyde
Vogel 3th.Edition 320p

[Edited on 30-7-2013 by Waffles SS]

Paddywhacker - 31-7-2013 at 02:28

In that case I should tell you that there is no way in the word that you will get an aldehyde by dehydrating butanol. You know butanol has just one oxygen, and water has one oxygen, and butyradehyde has one oxygen, so if you subtract water from butanol then the result won't have any oxygens left.

About your oxidation.... you do know that KMnO4 reduces to an alkaline solution... just look at the equation. So you will end up with a butyrate salt. In order to get any butyric acid you could possibly filter off the manganese dioxide, evaporate the filtrate to dryness, and then distill with a little H2SO4.

Hockeydemon - 31-7-2013 at 03:05

Quote: Originally posted by Paddywhacker

In that case I should tell you that there is no way in the word that you will get an aldehyde by dehydrating butanol. You know butanol has just one oxygen, and water has one oxygen, and butyradehyde has one oxygen, so if you subtract water from butanol then the result won't have any oxygens left.

Source

Quote:

When the vapours of a primary or a secondary alcohol are passed over heated copper at 573 K, dehydrogenation takes place and an aldehyde or a ketone is formed while tertiaryalcohols undergo dehydration.

Quote: Originally posted by Paddywhacker

About your oxidation.... you do know that KMnO4 reduces to an alkaline solution... just look at the equation. So you will end up with a butyrate salt. In order to get any butyric acid you could possibly filter off the manganese dioxide, evaporate the filtrate to dryness, and then distill with a little H2SO4.

Source

Am I missing something?

[Edited on 31-7-2013 by Hockeydemon]

adamsium - 31-7-2013 at 05:23

Quote: Originally posted by Paddywhacker

that there is no way in the word that you will get an aldehyde by dehydrating butanol

He is not trying to dehydrate anything. He is attempting an oxidation - specifically, a dehydrogenation - of butanol to butanal and, ultimately, to butyric acid.

I'd also submit that posts with no content other than, essentially, "this should be in beginnings" are entirely useless.

Waffles SS - 18-10-2013 at 11:37

Oxidation of Dibutyl Ether with Sodium Bromate lead to Butyl Butyrate and Butyric acid

Quote:

Autocatalytic Oxidation of Ethers with Sodium Bromate
Leonid Metsger, Shmuel Bittner
Tetrahedron 56 (3): 1905–10. doi:10.1016/S0040-4020(00)00098-3

[Edited on 19-10-2013 by Waffles SS]